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Frontier Molecular Orbital Theory

CHEM 430

Great Books on Frontier Orbital Theory

Basics of Frontier Molecular Orbital Theory

Since the majority of energy gain in a reaction between two molecules is a result of the HOMO of one molecule reacting with the LUMO of a second molecule this interaction is

called a Frontier Molecular Orbital (FMO) interaction

A reaction is thus favored when the HOMO (nucleophile) is unusually high in energy and the LUMO (electrophile) is unusually low in energy

What does unusually high HOMO or unusually low LUMO mean?

Must be compared relative to something -usually compare energy levels with a known unreactive C-H (or C-C) single bond

If the HOMO of a new compound is higher in energy than the HOMO of the C-H bond, then it will be more reactive as a nucleophile

If the LUMO of a new compound is lower in energy than the LUMO of the C-H bond, then it will be more reactive as an electrophile

How much higher or lower in energy will determine the relative rates of reactions

Frontier Molecular Orbitals – HOMO and LUMO

It makes sense that the HOMO and LUMO are the orbitals most likely to be involved in chemical reactivity. • Chemical reactions involve the redistribution of electrons (creation and

destruction of bonds, oxidation, reduction, …) • The HOMO is the orbital of highest energy that is still occupied, so

energetically it is the easiest to remove electrons from this orbital. This could be simply donating electron density to form a bond (act as a Lewis base) or it could be oxidation.

• The LUMO is the lowest lying orbital that is empty, so energetically it is the

easiest to add more electrons into this orbital…Lewis acid; reduction. • It isn’t always the HOMO and/or LUMO involved in chemical reactivity.

Symmetry plays a role, too. If the HOMO or LUMO isn’t of the correct symmetry, it might be the HOMO-1 or the LUMO+1 that is involved in the reaction.


Basics of Frontier Molecular Orbital TheoryWe can compare the placement of HOMO and LUMO levels

relative to placement of C-H bonds

C-H

*C-H

sp3C 1s H

A sp3 hybridized carbon atom and a 1s orbital of hydrogen have similar energy levels and strong overlap,

therefore high mixing

Factors that can adjust MO energy levels:1) Unmixed valence shell atomic orbitals

H+

No electrons in atomic orbital, therefore very electrophilic

:NH3

:OH2

Lone pair of electrons placed in atomic orbital

Because nitrogen is more electronegative than carbon,

orbital is lower in energy (likewise oxygen is lower than

nitrogen)

Both are very nucleophilic, ammonia more than water

Very low HOMO, therefore poor nucleophile

Very high LUMO, therefore poor electrophile



C-H

*C-H

sp3C 1s H




H+

:NH3

:OH2

2) Electric charge

OH

CH3

Negative charge will raise the energy of orbital, therefore make compound more nucleophilic





C-H

*C-H

sp3C 1s H




The degree of mixing of two orbitals is related to the amount of overlap between the orbitals

When two p orbitals overlap to form a bond, the orbitals begin higher in energy than a

hybridized orbital and the amount of overlap is less

2) Electric charge



3) Poor overlap of atomic orbitals

2p C 2p C

C-C

*C-C This makes HOMO into a good nucleophile



C-H

*C-H

sp3C 1s H




2) Electric charge



3) Poor overlap of atomic orbitals4) Poor energy match of orbitals

C-O

*C-O2p C

2p O

Since the oxygen 2p orbital is much lower in energy, the

energy match with carbon 2p is worse and therefore less mixing

This makes LUMO into a good electrophile


We can compare the placement of HOMO and LUMO levels relative to placement of C-H bonds

C-H

*C-H

sp3C 1s H




2) Electric charge



3) Poor overlap of atomic orbitals

sp3C

sp3Cl

C-Cl

*C-Cl

A C-Cl bond is good electrophile

4) Poor energy match of orbitals

C-Mg

*C-Mg

sp3C

spMg

A C-Mg bond is good nucleophile

Can also use orbital energy levels to understand differences in reactivity for C-X bonds

Basics of Frontier Molecular Orbital TheoryFrontier molecular orbital (FMO) theory allows a chemist to make predictions about a

reaction by knowing the placement of the HOMO and LUMO energy levels

A high HOMO level represents a compound that is a good nucleophile

A low LUMO level represents a compound that is a good electrophile

The energy level of the HOMO and LUMO can be predicted by knowing that when two atomic orbitals mix they form two new molecular orbitals,

one lower in energy and one higher in energyThe amount of mixing is dependent upon:

1) The amount of overlap between the mixing orbitals (e.g., the overlap for a bond is greater than the overlap for a bond)

2) The closer in energy are two orbitals, the greater the amount of mixing that occurs

OHCH3 NH2> >

C OR

RH3C Cl

Anything that will raise energy level of HOMO will

increase nucleophilicity

Anything that will lower energy level of LUMO will

increase electrophilicity


FMO will also allow prediction about where a reaction will occur (regiochemistry) and direction of approach (stereochemistry)

Consider a reaction with a carbonyl compound

FMO predicts that a carbonyl should react as an electrophile due to the low energy LUMO

The regio- and stereochemistry can also be predicted by considering the interacting frontier orbital (the LUMO)

rotate

LUMO of formaldehyde

The coefficient on carbon is larger than the coefficient on oxygen, therefore nucleophile

reacts at carbon


What direction should a nucleophile approach the carbonyl?

NUC

NUC

Expect this direction to be highly disfavored due to orthogonal

interaction with orbitals

Direction appears better, but still not optimal interaction

NUC

Optimal interaction (best overlap of

interacting orbitals)

Could there possibly be a method to test the angle of approach of nucleophile to carbonyl?

X-ray structures come to the solution once again!

Basics of Frontier Molecular Orbital TheoryWhat direction should a nucleophile approach the carbonyl?

NUC

Optimal interaction (best overlap of

interacting orbitals)

Could there possibly be a method to test the angle of approach of nucleophile to carbonyl?X-ray structures come to the solution once again!

Bürgi, H.B., Dunitz, J.D., Shefter,E., J. Am. Chem. Soc. (1973), 95, 5065-5067

Studied a variety of X-ray structures where a N

reacts with a carbonyl intramolecularly

As the N came closer to carbonyl, the C-O bond

lengthened and the carbonyl carbon becomes

pyramidalized

The angle of <N-C-O averaged 107˚ ( )

Called the “Bürgi-Dunitz” angle

Basics of Frontier Molecular Orbital TheoryWhat about the stereochemistry for a reaction with an alkyl halide?

Since alkyl halide is reacting as the electrophile, need to look at the LUMO

LUMO of methyl halide

Largest coefficient is on the backside of the carbon

NUC

So called “inversion of configuration”

Nucleophile thus reacts with a methyl halide in a SN2 reaction

with backside attack

LUMO of 2˚ alkyl halide

baseBonds that break

New bond

The base will abstract the hydrogen that is anticoplanar

to leaving group

Base thus reacts by abstracting hydrogen anticoplanar to

leaving group and form new bond in E2 reaction

General Reviews

Fleming, I. Frontier Orbitals and Organic Chemical Reactions

Fukui, Acc. Chem. Res. 1971, 4, 57.

Kirby, A. J. Stereoelectronic Effects.

+ Br:–

minor

major

Br: –Nu:

During the course of chemical reactions, the interaction of the highest filled (HOMO) and lowest unfilled (antibonding)

molecular orbital (LUMO) in reacting species is very important to the stabilization of the transition structure.

Geometrical constraints placed upon ground and transition statesby orbital overlap considerations.

Stereoelectronic Effects

Nonbonding interactions (van der Waals repulsion) between substituents within a molecule or between reacting molecules

Steric Effects

Universal Effects Governing all Chemical Reactions

Nondirectional Electronic Effects (Inductive Effects):

SN1

rate decreases as R becomes more electronegative

C Br

Me

RR

C RR

Me

Nu

R

H

R OOH

O

OH

R

R

HO

CR

R

Me

Br C MeR

R

Fukui Postulate for reactions:

SN2

General Reaction Types

Radical Reactions (~10%): A• B•+ A B

Polar Reactions (~90%): A(:) B(+)+ A B

Lewis Base Lewis Acid

FMO concepts extend the donor-acceptor paradigm to non-obvious families of reactions

QAQB

Q: Charge densityε: Dielectric constantR: distance of A to Bc: coefficient of MO m of species A, or MO n of species Bβ: Overlap IntegralE: Energy of MO

∆E =(cm

AcnBβ)2

(Em - En)mnεR

Coulomb Term Orbital Term

Consider stabilization energy (∆E) when bringing atoms A & B together:

2ΣΣ

Steric, Inductive & Stereoelectronic Effects

The H2 Molecule (again!!)Let's combine two hydrogen atoms to form the hydrogen molecule.Mathematically, linear combinations of the 2 atomic 1s states createtwo new orbitals, one is bonding, and one antibonding:

Ener

gy

1s 1s

σ∗ (antibonding)

Rule one: A linear combination of n atomic states will create n MOs.

∆E

∆E

Let's now add the two electrons to the new MO, one from each H atom:

Note that ∆E1 is greater than ∆E2. Why?

σ (bonding)

σ (bonding)

∆E2

∆E1

σ∗ (antibonding)

1s1s

ψ2

ψ2

ψ1

ψ1

Ener

gy

+C1ψ1σ = C2ψ2

Linear Combination of Atomic Orbitals (LCAO): Orbital Coefficients

Each MO is constructed by taking a linear combination of the individual atomic orbitals (AO):

Bonding MO

Antibonding MO C*2ψ2σ∗ =C*1ψ1–

The coefficients, C1 and C2, represent the contribution of each AO.

Rule Two: (C1)2 + (C2)2 = 1

= 1antibonding(C*1)2+bonding(C1)2 Rule Three:

Ener

gyπ∗ (antibonding)

π (bonding)

Consider the pi-bond of a C=O function: In the ground state pi-C–Ois polarized toward oxygen. Note (Rule 2) that the antibonding MOis polarized in the opposite direction.

C

C

O

C O

The squares of the C-values are a measure of the electron populationin neighborhood of atoms in question

In LCAO method, both wave functions must each contribute one net orbital

Dihydrogen Molecular Orbitals

Weak bonds will have corresponding low-lying antibonds.

π Si–Si = 23 kcal/molπ C–Si = 36 kcal/molπ C–C = 65 kcal/molThis trend is even more dramatic with pi-bonds:

σ∗ C–Siσ∗ C–C

σ C–Si

σ C–C

Bond length = 1.87 ÅBond length = 1.534 ÅH3C–SiH3 BDE ~ 70 kcal/molH3C–CH3 BDE = 88 kcal/mol

The following generalizations on covalent bonding are useful.

When one compares bond strengths between C–C and C–X, where X is some other element such as O, N, F, Si, or S, keep in mind that covalent and ionic contributions vary independently. Hence, the

mapping of trends is not a trivial exercise.

Bond Energy (BDE) = δ Ecovalent + δ Eionic

Bond strengths (Bond dissociation energies) are composed of a covalent contribution (δ Ecov) and an ionic contribution (δ Eionic).

better than

For example, consider elements in Group IV, Carbon and Silicon. We know that C-C bonds are considerably stronger by Ca. 20 kcal mol-1

than C-Si bonds.

Overlap between orbitals of comparable energy is more effective than overlap between orbitals of differing energy.

Formation of a weak bond will lead to a corresponding low-lying antibonding orbital. Such structures are reactive as both nucleophiles & electrophiles

••

Better than

Better than

Case-2: Two anti sigma bonds

σ C–YHOMO

σ* C–XLUMO

σ* C–XLUMO

lone pairHOMO

σ* C–XLUMO

σ* C–XLUMO

lone pairHOMO

Case-1: Anti Nonbonding electron pair & C–X bond

An anti orientation of filled and unfilled orbitals leads to better overlap.

This is a corrollary to the preceding generalization. There are two common situations.

Better than

For π Bonds:

For σ Bonds:

Orbital orientation strongly affects the strength of the resulting bond.

Better than

This is a simple notion with very important consequences. It surfaces inthe delocalized bonding which occurs in the competing anti (favored) syn (disfavored) E2 elimination reactions. Review this situation.

σ C–YHOMO

A C A C

C CC C

A C

X

A

Y

C

X

A B A B

C-SP3

Si-SP3

C-SP3C-SP3

SiC Si CCCCC

Y

Y

X X

XX

BABA

Bonding Generalizations

σ∗Csp3-Csp2 is a better acceptor orbital than σ∗Csp3-Csp3

C-sp3

C-sp3

σ* C-C

σ C-C

C-sp3

σ C-C

σ* C-C

C-sp2

Donor Acceptor Properties of Csp3-Csp3 & Csp3-Csp2 Bonds

The greater electronegativity of Csp2 lowers both the bonding & antibonding C–C states. Hence:

σ Csp3-Csp3 is a better donor orbital than σ Csp3-Csp2

σ∗C-O is a better acceptor orbital than σ∗C-C

σ C-C is a better donor orbital than σ C-O

The greater electronegativity of oxygen lowers both the bonding & antibonding C-O states. Hence:

Consider the energy level diagrams for both bonding & antibonding orbitals for C-C and C-O bonds.

Donor Acceptor Properties of C-C & C-O Bonds

O-sp3

σ* C-O

σ C-O

C-sp3

σ C-C

σ* C-C

better donor

better acceptor

decreasing donor capacity

Nonbonding States

poorest donor

The following are trends for the energy levels of nonbonding states of several common molecules. Trend was established by

photoelectron spectroscopy.

best acceptor

poorest donor

Increasing σ∗-acceptor capacity

σ-anti-bonding States: (C–X)

σ-bonding States: (C–X)

decreasing σ-donor capacity

Following trends are made on the basis of comparing the bonding and antibonding states for the molecule CH3-X where X = C, N, O, F, & H.

Hierarchy of Donor & Acceptor States

CH3–CH3

CH3–H

CH3–NH2

CH3–OH

CH3–F

CH3–H

CH3–CH3

CH3–NH2

CH3–OH

CH3–F

HCl:H2O:

H3N:H2S:

H3P:

Donor-Acceptor Properties of Bonding & Antibonding States

3p Orbital

This becomes apparent when the radial probability functions for s and p-states are examined: The radial probability functions for the

hydrogen atom s & p states are shown below.

3s Orbital

Electrons in 2s states "see" a greater effective nuclear charge than electrons in 2p states. This correctly implies that the stability of nonbonding electron

pairs is directly proportional to the % of s-character in the doubly occupied orbital

Least stable Most stable

The above trend indicates that the greater the % of s-character at a given atom, the greater the electronegativity of that atom.

There is a direct relationship between %s character & hydrocarbon acidity

There is a linear relationship between %s character & Pauling electronegativity

Å

Rad

ial P

roba

bilit

y100 %

2p Orbital

2s Orbital2s Orbital

1s Orbital

100 %

Rad

ial P

roba

bilit

y

Å

Electrons in s-states "see" a higher nuclear charge. This is even more obvious in an electron density map (see

http://www.shef.ac.uk/chemistry/orbitron/). The s-orbitals have maximal electron density at the nucleus, and the p-orbitals have none.

Csp3 Csp2 Csp

Hybridization vs. Electronegativity

[F5Sb–F–SbF5]–

The Adamantane Reference(MM2)

T. Laube, Angew. Chem. Int. Ed. 1986, 25, 349

First X-ray Structure of an Aliphatic Carbocation

110 °100.6 °

1.530 Å1.608 Å

1.528 Å1.431 Å

Bonds participating in the hyperconjugative interaction, e.g. C–R, will be lengthened while the C(+)–C bond will be shortened.

Physical Evidence for Hyperconjugation

The new occupied bonding orbital is lower in energy. When you stabilize the electrons in a system you stabilize the system itself.

Take a linear combination of σ C–R and Csp2 p-orbital:

σ C–R

σ∗ C–R

σ C–R

σ∗ C–R

The Molecular Orbital Description

coplanar orientation between interacting orbitalsStereoelectronic Requirement for Hyperconjugation:

The graphic illustrates the fact that the C-R bonding electrons can "delocalize" to stabilize the electron deficient carbocationic center.

Note that the general rules of drawing resonance structures still hold:the positions of all atoms must not be changed.

+

The interaction of a vicinal bonding orbital with a p-orbital is referred to as hyperconjugation.

Me

Me

Me

H

C C

R

H

H H

HC

H

HCH

H

CH

HC

H

H

Me

Me

Me

C

R

This is a traditional vehicle for using valence bond to denote charge delocalization.

+

+

C

H

H H

H

bonding interaction

More substituted carbocations have more adjacent C-R bonds to act as donors to the empty p orbital

Hence, more substituted carbocations are more stable.

C C

H

H

H H

H+

H CH

H+

C C

H

H

H H

CH3

+C C

H

H

H CH3

CH3

+

least stable most stable

Me 1o 2o 3o

Hyperconjugation: Carbocation Stabilization

NMR Spectroscopy Greater e-density at R

Less e-density at X NMR Spectroscopy

Longer C–R bond X-ray crystallography

Infrared Spectroscopy Weaker C–R bond

Stronger C–X bond Infrared Spectroscopy

X-ray crystallography Shorter C–X bond

Spectroscopic ProbeChange in Structure

The Expected Structural Perturbations

As the antibonding C–R orbital decreases in energy, the magnitude

of this interaction will increase σ C–R

σ∗ C–R

The Molecular Orbital Description

Delocalization of nonbonding electron pairs into vicinal antibonding orbitals is also possible

X

Since nonbonding electrons prefer hybrid orbitals rather than p orbitals, this orbital can adopt either a syn or anti relationship to

the vicinal C–R bond.

C X

R

HH

HH X H

HCH

H

R–

This delocalization is referred to as "negative" hyperconjugation antibonding σ∗ C–R

Overlap between two orbitals is better in the anti orientation as stated in "Bonding Generalizations" handout.

+

–

Anti Orientation

filled hybrid orbital

filled hybrid orbital

antibonding σ∗ C–RSyn Orientation

–

+C X

HH

C X

HHCH

CHH

R

X

H

R

XC X

HH

C X

HH

R:

R:

Nonbonding e– pair

–

R

R

"Negative" Hyperconjugation & Anomeric Effects

We now conclude that this is another example of negative hyperconjugation.

filled N-sp2

The low-frequency shift of the cis isomer is a result of N–H bond weakening due to the anti lone pair on the adjacent (vicinal)

nitrogen which is interacting with the N–H antibonding orbital. Note that the orbital overlap is not nearly as good from the trans isomer.

Infrared evidence for lone pair delocalization into vicinal antibonding orbitals.

ν N–H = 2188 cm -1

ν N–H = 2317 cm -1

filled N-sp2

antibonding σ∗ N–H

antibonding σ∗ N–H

The N–H stretching frequency of cis-methyl diazene is 200 cm-1 lower than the trans isomer.

N. C. Craig & co-workers JACS 1979, 101, 2480.

ν C–H = 3050 cm -1ν C–H = 2730 cm -1

Aldehyde C–H Infrared Stretching Frequencies

The IR C–H stretching frequency for aldehydes is lower than the closely related olefin C–H stretching frequency. For years this

observation has gone unexplained.

The Anomeric Effect

It is not unexpected that the methoxyl substituent on a cyclohexane ring prefers to adopt the equatorial conformation.

∆ G° = +0.8 kcal/mol

∆ G° = –0.6 kcal/mol

What is unexpected is that the closely related 2-methoxytetrahydropyranprefers the axial conformation:

That effect which provides the stabilization of the axial OR conformerwhich overrides the inherent steric bias of the substituent is referred to asthe anomeric effect.

axial O lone pair↔σ∗ C–H axial O lone pair↔σ∗ C–Opreferred

Principal HOMO-LUMO interaction from each conformation is illustrated below:

Since the antibonding C–O orbital is a better acceptor orbital than the antibonding C–H bond, the axial OMe conformer is better stabilized by

this interaction which is worth ca 1.2 kcal/mol.Other electronegative substituents such as Cl, SR etc. also participate in

anomeric stabilization.

This conformer preferred by 1.8 kcal/mol

1.819 Å

1.781 Å

Why is axial C–Cl bond longer ?

N N

Me H

N

H

N

Me

CH

C

RO

HC

R R

R

N NMe

H

OMe H

OMe

OMe

H

N N

Me

OMe

H

OO

O

H

OMe O H

OMe

Cl

HO O O

H

Cl H

Cl

H

H

The Anomeric Effect & Related Issues

Klopman-Salem Equation & FMO Theory

“For a research worker, the unforgotten moments of life are those rare ones, which come after years of plodding work, when the veil over nature s secret seems suddenly to lift and when what was dark and chaotic appears in a clear and beautiful light and pattern.”

-Gerty Cori, Washington U, Nobel Laureate, 1947

Sterics Electrostatics HOMO-LUMO

Klopman-Salem EquationDescribes the energetic changes that occurs when two speciesapproach each other - as they begin to interact their MO’s overlapand atoms bearing partial charges experience electrostatic forces

E

E

EE

CH2=CH2

+1.5 eV

–10.52 eV

+– –

––

– –

–

–

+

++

+

+

+ +

OMe

OMe

CO2Me

CO2Me 0 eV

–10.72 eV

+0.71 –0.71

+0.71 +0.71

+0.66

+0.61

+1.5 eV+2.0 eV

–10.52 eV–9.05 eV

CH2=CHOMe

+0.39

–0.72+0.69–0.47

+0.43 +0.33

CH2=CH2CH2=CHCO2Me

HOMO and LUMO Energies and Orbital Coefficients of Common Alkenes

X X

LUMO

c1 c2 c3 c4

HOMO

Alkene HOMO (eV) c1 c2 LUMO (eV) c3 c4

CH2=CH2a –10.52 0.71 0.71 +1.5 0.71 –0.71

CH2=CHCla –10.15 0.44 0.30 +0.5 0.67 –0.54

CH2=CHMea –9.88 0.67 0.56 +1.8a,b 0.67 –0.65

MeCH=CHMe –9.13c +2.22d

EtCH=CH2 –9.63e +2.01e

–8.94f,g +2.1g

CH2=CHOMea –9.05;–8.93c 0.61 0.39 +2.0 0.66 –0.72

CH2=CHSMea –8.45 0.34 0.17 +1.0 0.63 –0.48

CH2=CHNMe2a –9.0 0.50 0.20 +2.5 0.62 –0.69

CH2=CHCO2Me –10.72 0.43 0.33 0 0.69 –0.47

CH2=CHCN –10.92 0.60 0.49 0 0.66 –0.54

CH2=CHNO2a –11.4 0.62 0.60 +0.7 0.54 –0.32

CH2=CHPh –8.48 0.49 0.32 +0.8 0.48 –0.33

CH2=CHCHO –10.89b 0.58a 0.48a +0.60b 0.404b –0.581b

CH2=CHCHO/BF3b –12.49 +0.43 0.253 –0.529

CH2=CHCO2Hb –10.93 +2.91 0.461 –0.631

O O–10.29h –1.91i

Ph

PhCH=CH2

CH=CH2

–0.42

CH2=CH2CH2=CH—CH=CH2

+0.41 –9.03 eV+0.57

+0.56 +0.71 –0.71

+0.71 +0.71

+0.48

+0.49

+1.5 eV+0.8 eV

–10.52 eV

–8.48 eV+0.32

–0.33

CH2=CHPh

+1.0 eV

+ 1.0 eV

HOMO

LUMO

+0.560–0.42

–0.57–0.41+0.41

+0.57

–0.42+0.56

– 9.03 eV

Z

Z

Z

Z

X

X

X

X

– 8.7 eV

+ 2.3 eV

– 8.5 eV– 9.3 eV– 9.5 eV

+ 2.5 eV

– 0.3 eV– 0.5 eV

– 9.1 eV

+ 1.0 eV

HOMO, LUMO Energies and Orbital Coefficients for Substituted Dienes

R1 R1

R2R2

c4c2c1

LUMOHOMO

c3

D i e n e HOMOa c1 c2 LUMO c3 c4

a

–9.07;–8.85b 0.57 –0.57 +1.0;3.38b 0.56 0.56

Meb

–9.78a;–8.54 0.314 0.315 +3.51 0.629c 0.617c

Me b

–9.04a;–8.72 0.340 0.296 +3.38 0.56d 0.55d

Me Me

–8.76a 2.18e

Me Me –8.39a

Ph –8.16a 0.408f 0.416f

Ph

–8.77a 0.572g 0.335g

MeO –8.21a;–8.24b 0.235b 0.313b +3.77b 0.644c 0.609c

OMe

–8.62a 0.352b 0.103b +3.60b

SMeb

–7.94 0.240 0.256 +3.25

SMe

b –8.37 0.399 0.201 +3.25

LUMOalkene

HOMOdiene

+0.71

–0.71+0.71+1.5 eV

–10.52 eV

+0.56–0.42

+0.57+0.41 –0.41

–0.57

–0.42+0.56

+ 1.0 eV

–9.03 eV

+0.71

–9.05 eV

–10.72 eV–10.52 eV

+2.0 eV

0 eV

+1.5 eV

+1.0 eV

–9.07 eV

E = 11.07 eV

E = 10.57 eV

E = 9.07 eV

MeOEtO2C

HOMOdiene – LUMOalkene Gap for Butadiene & Substituted Alkenes

Regioselectivity of 1-Alkyl-1,3-butadienes with Electron Deficient Alkenes

ZRR

Z

R

Z

+ +

A B

R Z A / B

Me CHO 8:1

Me C N 10:1

Me CO2Me 6.8:1

i-Pr CO2Me 5:1

n-Bu CO2Me 5.1:1

t-Bu CO2Me 4.1:1

CO2Me

HPh

Ph

CO2Me

H

–0.475

0.625–0.47

0.69

O

OMePh

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